A LEVEL CHEMISTRY
TOPIC 12 – ACIDS, BASES AND BUFFERS
ASSESSED HOMEWORK
Answer all questions
Max 80 marks

Name		……………………………………………………………..

Mark		……../80	……....%	Grade ………

1.	
A student was given a task to determine the percentage purity of a sample of salicylic acid. The method used by the student to prepare a solution of salicylic acid is described below.
•        0.500 g of an impure sample of salicylic acid was placed in a weighing bottle.
•        The contents were tipped into a beaker and 100 cm3 of distilled water were added.
•        Salicylic acid does not dissolve well in cold water so the beaker and its contents were heated gently until all the solid had dissolved.
•        The solution was poured into a 250 cm3 graduated flask and made up to the mark with distilled water.
(a)     Give two additional instructions that would improve this method for making up the salicylic acid solution.
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(2)
(b)     The pH of this solution was measured and a value of 2.50 was obtained. 
Calculate the concentration of salicylic acid in this solution. 
Assume that salicylic acid is the only acid in this solution. The Ka for salicylic acid is 1.07 × 10–3 mol dm–3. You may represent salicylic acid as HA. 
Show your working.
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(3)
(c)     Use your answer to part (b) to calculate the mass of salicylic acid (Mr = 138.0) present in the original sample. 
(If you were unable to complete the calculation in part (b), assume that the concentration of salicylic acid is 8.50 × 10–3 mol dm–3. This is not the correct answer.)
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(2)

(d)     Use your answer to part (c) to calculate the percentage purity of the salicylic acid used to make the solution. 
(If you were unable to complete the calculation in part (c), assume that the mass of salicylic acid is 0.347 g. This is not the correct answer.)
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(1)
(Total 8 marks)

2.	In this question, give all values of pH to two decimal places.
Calculating the pH of aqueous solutions can involve the use of equilibrium constants such as Kw and Ka
Kw is the ionic product of water. The value of Kw is 5.48 × 10–14 mol2 dm–6 at 50 °C.
(a)     (i)      Write an expression for pH.
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(1)
(ii)     Write an expression for Kw
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(1)
(b)     (i)      Calculate the pH of pure water at 50 °C.
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(2)
(ii)     Suggest why this pure water is not acidic.
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(1)
(iii)     Calculate the pH of 0.140 mol dm–3 aqueous sodium hydroxide at 50 °C.
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(3)

(c)     	Calculate the pH of the solution formed when 25.0 cm3 of 0.150 mol dm–3 aqueous sulfuric acid 
	are added to 30.0 cm3 of 0.200 mol dm–3 aqueous potassium hydroxide at 25 °C. Assume that 	the sulfuric acid is fully dissociated.
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(6)
(Total 14 marks)

3.	This question is about several Brønsted–Lowry acids and bases.
(a)     Define the term Brønsted–Lowry acid.
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(1)
(b)     Three equilibria are shown below. For each reaction, indicate whether the substance immediately above the box is acting as a Brønsted–Lowry acid (A) or a Brønsted–Lowry base (B) by writing A or B in each of the six boxes.

 (i)	CH3COOH	+	H2O
		CH3COO–
	+	H3O+	

(1)
 (ii)	CH3NH2	+	H2O
		CH3NH3+
	+	OH–	

(1)
(iii)	   HNO3
	+	H2SO4		H2NO3+
	+	HSO4 –	

(1)
(c)     A 25.0 cm3 sample of 0.0850 mol dm–3 hydrochloric acid was placed in a beaker.
Distilled water was added until the pH of the solution was 1.25.
Calculate the total volume of the solution formed. State the units.
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 (3)

(d)     At 298 K, the value of the acid dissociation constant (Ka) for the weak acid HX in aqueous solution is 3.01 × 10–5 mol dm–3.
(i)      Calculate the value of pKa for HX at this temperature.
Give your answer to 2 decimal places.
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(1)
(ii)     Write an expression for the acid dissociation constant (Ka) for the weak acid HX.
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(1)
(iii)     Calculate the pH of a 0.174 mol dm–3 solution of HX at this temperature.
Give your answer to 2 decimal places.
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 (3)
(e)     An acidic buffer solution is formed when 10.0 cm3 of 0.125 mol dm–3 aqueous sodium hydroxide are added to 15.0 cm3 of 0.174 mol dm–3 aqueous HX.
The value of Ka for the weak acid HX is 3.01 × 10–5 mol dm–3.
Calculate the pH of this buffer solution at 298 K.
Give your answer to 2 decimal places.
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(6)
(Total 18 marks)
 4.     A 0.210 mol dm–3 solution of potassium hydroxide was added from a burette to 25.0 cm3 of a 0.160 mol dm–3 solution of ethanoic acid in a conical flask.
Given that the value of the acid dissociation constant, Ka, for ethanoic acid is 1.74 × 10–5 mol dm–3, calculate the pH at 25 °C of the solution in the conical flask at the following three points:
before any potassium hydroxide had been added;
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after 8.0 cm3 of potassium hydroxide solution had been added;
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after 40.0 cm3 of potassium hydroxide solution had been added.
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(Total 16 marks)

 

5.	The acid dissociation constant, Ka, for ethanoic acid is given by the expression
Ka = 
The value of Ka for ethanoic acid is 1.74 × 10−5 mol dm−3 at 25 °C.
(a)     A buffer solution is prepared using ethanoic acid and sodium ethanoate. In the buffer solution, the concentration of ethanoic acid is 0.186 mol dm−3 and the concentration of sodium ethanoate is 0.105 mol dm−3.
Calculate the pH of this buffer solution.
Give your answer to 2 decimal places.
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 (3)
(b)     In a different buffer solution, the concentration of ethanoic acid is 0.251 mol dm−3 and the concentration of sodium ethanoate is 0.140 mol dm−3.
A sample of hydrochloric acid containing 0.015 mol of HCl is added to 1000 cm3 of this buffer solution.
Calculate the pH of the buffer solution after the hydrochloric acid has been added.
You should ignore any change in total volume.
Give your answer to 2 decimal places.
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 (5)
(Total 8 marks)

6.	Titration curves, labelled E, F, G and H, for combinations of different aqueous solutions of 	acids and bases are shown below.
All solutions have concentrations of 0.1 mol dm−3.

(a)     In this part of the question, write the appropriate letter in each box.
From the curves E, F, G and H, choose the curve produced by the addition of
 
 	(i)	sodium hydroxide to 25 cm3 of ethanoic acid		
(1)
 	(ii)	ammonia to 25 cm3 hydrobromic acid		
(1)
 	(iii)	hydrochloric acid to 25 cm3 of potassium hydroxide		
(1)

(b)     The table shows information about some acid-base indicators.
 
 	Indicator	pH range	Lower pH colour	Higher pH colour	
 	pentamethoxy red	1.2−3.2	violet	colourless	
 	naphthyl red	3.7−5.0	red	yellow	
 	4−nitrophenol	5.6−7.0	colourless	yellow	
 	cresol purple	7.6−9.2	yellow	purple	

	(i)      	Which indicator in the table could be used for the titration that produces curve E but 
		not for the titration that produces curve F?
Tick (✓) one box.
 	pentamethoxy red		
 	naphthyl red		
 	4−nitrophenol		
 	cresol purple		
(1)
(ii)     Give the colour change at the end point of the titration that produces curve H when naphthyl red is used as the indicator.
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(1)
(iii)    A beaker contains 25 cm3 of a buffer solution at pH = 6.0
Two drops of each of the four indicators in the table are added to this solution.
State the colour of the mixture of indicators in this buffer solution.
You should assume that the indicators do not react with each other.
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(1)
(Total 6 marks)

7.	When iron(II) sulfate is used for killing weeds in lawns, it is often mixed with the fertiliser ammonium sulfate. Ammonium sulfate also makes the soil acidic.
(a)     Write an equation to show how the ammonium ion behaves as a Brønsted–Lowry acid in water.
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(1)
(b)     Compounds such as ammonium sulfate react on warming with sodium hydroxide solution as shown in the equation below.
(NH4)2SO4  +  2NaOH   Na2SO4  +  2NH3  +  2H2O
Use this information to describe a simple test, other than smell, to show that ammonia is evolved. State what you would observe.
Test ................................................................................................................
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Observation ...................................................................................................
(2)
(Total 3 marks)

8.	Describe briefly how you would ensure that a reading from a pH meter is accurate.
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(Total 2 marks)

9.	This question is about the following reaction scheme which shows the preparation of polymer P.
 
K is a weak acid with a pKa of 9.95. The pH of a 0.10 mol dm−3 solution of K is
A       4.48
B       4.98
C       5.48
D       5.98
(Total 1 mark)

10.	A solution of sodium ethanoate has a pH of 8.91 at 25 °C. The hydrogen ion and hydroxide ion concentrations in this solution are
A       [H+] = 1.00 × 10−9 mol dm−3 [OH−] = 1.00 × 10−5 mol dm−3
B       [H+] = 1.00 × 10−9 mol dm−3 [OH−] = 8.13 × 10−6 mol dm−3
C       [H+] = 1.23 × 10−9 mol dm−3 [OH−] = 1.00 × 10−5 mol dm−3
D       [H+] = 1.23 × 10−9 mol dm−3 [OH−] = 8.13 × 10−6 mol dm−3
(Total 1 mark)

11.	Which one of the following is the change in units of pH which occurs when 10.0 cm3 of a 1.0 M 	solution of a strong monoprotic acid are made up to 1.0 dm3 with water?
A       1
B       2
C       3
D       5
(Total 1 mark)

12.	Which one of the following could be true in an aqueous solution of sodium hydroxide?
A       [H+] = [OH−]
B       pH = −log10 [OH−]
C       pH = 1.2
D       pH = 12.8
(Total 1 mark)

13.	Use the information below to answer this question.
A saturated solution of magnesium hydroxide, Mg(OH)2, contains 0.1166 g of Mg(OH)2 in 10.00 dm3 of solution. In this solution the magnesium hydroxide is fully dissociated into ions.
Which one of the following is the pH of a solution of magnesium hydroxide containing 4.0 × 10−5 mol dm−3 of hydroxide ions at 298 K? 
(Kw = 1.0 × 10−14 mol2 dm−6 at 298 K)
A       9.6
B       9.5
C       8.6
D       8.3
(Total 1 mark)